Why do we need the equation of state of an ideal gas? Equation of state of gases. Ideal gas mixture

DEFINITION

In order to make formulas and laws in physics easier to understand and use, various types of models and simplifications are used. Such a model is ideal gas. A model in science is a simplified copy of a real system.

The model reflects the most essential characteristics and properties of processes and phenomena. The ideal gas model takes into account only the basic properties of molecules that are required to explain the basic behavior of the gas. An ideal gas resembles a real gas in a fairly narrow range of pressures (p) and temperatures (T).

The most important simplification of an ideal gas is that the kinetic energy of the molecules is considered to be much greater than the potential energy of their interaction. Collisions of gas molecules are described using the laws of elastic collision of balls. Molecules are considered to move in a straight line between collisions. These assumptions make it possible to obtain special equations, which are called the equations of state of an ideal gas. These equations can be applied to describe the states of real gas at low temperatures and pressures. The equations of state can be called formulas for an ideal gas. We also present other basic formulas that are used in studying the behavior and properties of an ideal gas.

Equations of ideal state

Mendeleev-Clapeyron equation

where p is gas pressure; V is the volume of gas; T is the gas temperature on the Kelvin scale; m is gas mass; - molar mass of gas; — universal gas constant.

The equation of state of an ideal gas is also the expression:

where n is the concentration of gas molecules in the volume under consideration; .

Basic equation of molecular kinetic theory

Using a model such as an ideal gas, the basic equation of molecular kinetic theory (MKT) (3) is obtained. Which suggests that the pressure of a gas is the result of a huge number of impacts of its molecules on the walls of the vessel in which the gas is located.

where is the average kinetic energy of the translational motion of gas molecules; - concentration of gas molecules (N - number of gas molecules in the vessel; V - volume of the vessel); - mass of a gas molecule; - root mean square speed of the molecule.

Internal energy of an ideal gas

Since in an ideal gas the potential energy of interaction between molecules is assumed to be zero, the internal energy is equal to the sum of the kinetic energies of the molecules:

where i is the number of degrees of freedom of an ideal gas molecule; - Avogadro's number; - amount of substance. The internal energy of an ideal gas is determined by its thermodynamic temperature (T) and is proportional to its mass.

Ideal gas work

For an ideal gas in an isobaric process (), the work is calculated using the formula:

In an isochoric process, the work done by the gas is zero, since there is no change in volume:

For an isothermal process ():

For an adiabatic process () the work is equal to:

where i is the number of degrees of freedom of a gas molecule.

Examples of solving problems on the topic “Ideal gas”

EXAMPLE 1

Ideal gas is a gas in which there are no forces of mutual attraction and repulsion between molecules and the sizes of the molecules are neglected. All real gases at high temperatures and low pressures can practically be considered ideal gases.
The equation of state for both ideal and real gases is described by three parameters according to equation (1.7).
The equation of state of an ideal gas can be derived from molecular kinetic theory or from a joint consideration of the Boyle-Mariotte and Gay-Lussac laws.
This equation was derived in 1834 by the French physicist Clapeyron and for 1 kg of gas mass has the form:

Р·υ = R·Т, (2.10)

where: R is the gas constant and represents the work done by 1 kg of gas in a process at constant pressure and with a temperature change of 1 degree.
Equation (2.7) is called t thermal equation of state or characteristic equation .
For an arbitrary amount of gas of mass m, the equation of state will be:

Р·V = m·R·Т. (2.11)

In 1874, D.I. Mendeleev, based on Dalton’s law ( “Equal volumes of different ideal gases at the same temperatures and pressures contain the same number of molecules.”) proposed a universal equation of state for 1 kg of gas, which is called Clapeyron-Mendeleev equation:

Р·υ = R μ ·Т/μ , (2.12)

where: μ - molar (molecular) mass of gas, (kg/kmol);

R μ = 8314.20 J/kmol (8.3142 kJ/kmol) - universal gas constant and represents the work done by 1 kmol of an ideal gas in a process at constant pressure and with a temperature change of 1 degree.
Knowing R μ, you can find the gas constant R = R μ / μ.
For an arbitrary mass of gas, the Clapeyron-Mendeleev equation will have the form:

Р·V = m·R μ ·Т/μ. (2.13)

A mixture of ideal gases.

Gas mixture refers to a mixture of individual gases that enter into any chemical reactions with each other. Each gas (component) in the mixture, regardless of other gases, completely retains all its properties and behaves as if it alone occupied the entire volume of the mixture.
Partial pressure- this is the pressure that each gas included in the mixture would have if this gas were alone in the same quantity, in the same volume and at the same temperature as in the mixture.
The gas mixture obeys Dalton's law:
║The total pressure of the gas mixture is equal to the sum of the partial pressures║individual gases that make up the mixture.

P = P 1 + P 2 + P 3 + . . . Р n = ∑ Р i , (2.14)

where P 1, P 2, P 3. . . Р n – partial pressures.
The composition of the mixture is specified by volume, mass and mole fractions, which are determined respectively using the following formulas:

r 1 = V 1 / V cm; r 2 = V 2 / V cm; … r n = V n / V cm, (2.15)
g 1 = m 1 / m cm; g 2 = m 2 / m cm; … g n = m n / m cm, (2.16)
r 1 ′ = ν 1 / ν cm; r 2 ′ = ν 2 / ν cm; … r n ′ = ν n / ν cm, (2.17)

where V 1; V 2 ; … V n ; V cm – volumes of components and mixture;
m 1; m2; … m n ; m cm – masses of components and mixture;
ν 1; ν 2 ; … ν n ; ν cm – amount of substance (kilomoles)
components and mixtures.
For an ideal gas, according to Dalton's law:

r 1 = r 1 ′; r 2 = r 2 ′; … r n = r n ′ . (2.18)

Since V 1 +V 2 + … + V n = V cm and m 1 + m 2 + … + m n = m cm,

then r 1 + r 2 + … + r n = 1, (2.19)
g 1 + g 2 + … + g n = 1. (2.20)

The relationship between volume and mass fractions is as follows:

g 1 = r 1 ∙μ 1 /μ cm; g 2 = r 2 ∙μ 2 /μ cm; … g n = r n ∙μ n /μ cm, (2.21)

where: μ 1, μ 2, ... μ n, μ cm – molecular weights of the components and mixture.
Molecular weight of the mixture:

μ cm = μ 1 r 1 + r 2 μ 2 + … + r n μ n. (2.22)

Gas constant of mixture:

R cm = g 1 R 1 + g 2 R 2 + … + g n R n =
= R μ (g 1 /μ 1 + g 2 /μ 2 + … + g n /μ n) =
= 1 / (r 1 /R 1 + r 2 /R 2 + ... + r n /R n) . (2.23)

Specific mass heat capacities of the mixture:

with р cm. = g 1 with р 1 + g 2 with р 2 + … + g n with р n. (2.24)
with v see = g 1 with p 1 + g 2 with v 2 + ... + g n with v n. (2.25)

Specific molar (molecular) heat capacities of the mixture:

with rμ cm. = r 1 with rμ 1 + r 2 with rμ 2 + … + r n with rμ n. (2.26)
with vμ cm. = r 1 with vμ 1 + r 2 with vμ 2 + … + r n with vμ n. (2.27)

Topic 3. Second law of thermodynamics.

Basic provisions of the second law of thermodynamics.

The first law of thermodynamics states that heat can be converted into work, and work into heat, and does not establish the conditions under which these transformations are possible.
The transformation of work into heat always occurs completely and unconditionally. The reverse process of converting heat into work during its continuous transition is possible only under certain conditions and not completely. Heat can naturally move from hotter bodies to colder ones. The transfer of heat from cold bodies to heated ones does not occur by itself. This requires additional energy.
Thus, for a complete analysis of phenomena and processes, it is necessary to have, in addition to the first law of thermodynamics, an additional law. This law is second law of thermodynamics . It establishes whether a particular process is possible or impossible, in which direction the process proceeds, when thermodynamic equilibrium is achieved, and under what conditions maximum work can be obtained.
Formulations of the second law of thermodynamics.
For the existence of a heat engine, 2 sources are needed - hot spring and cold spring (environment). If a heat engine operates from only one source, it is called perpetual motion machine of the 2nd kind.
1 formulation (Ostwald):
| "A perpetual motion machine of the 2nd kind is impossible."

A perpetual motion machine of the 1st kind is a heat engine in which L>Q 1, where Q 1 is the supplied heat. The first law of thermodynamics “allows” the possibility of creating a heat engine that completely converts the supplied heat Q 1 into work L, i.e. L = Q 1. The second law imposes more stringent restrictions and states that the work must be less than the heat supplied (L A perpetual motion machine of the 2nd kind can be realized if heat Q 2 is transferred from a cold source to a hot one. But for this, heat must spontaneously transfer from a cold body to a hot one, which is impossible. This leads to the 2nd formulation (by Clausius):
|| "Heat cannot spontaneously transfer from more
|| cold body to a warmer one."
To operate a heat engine, two sources are needed - hot and cold. 3rd formulation (Carnot):
|| "Where there is a temperature difference, it is possible to commit
|| work."
All these formulations are interconnected; from one formulation you can get another.

Entropy.

One of the functions of the state of a thermodynamic system is entropy. Entropy is a quantity defined by the expression:

dS = dQ / T. [J/K] (3.1)

or for specific entropy:

ds = dq / T. [J/(kg K)] (3.2)

Entropy is an unambiguous function of the state of a body, taking on a very specific value for each state. It is an extensive (depending on the mass of the substance) state parameter and in any thermodynamic process is completely determined by the initial and final state of the body and does not depend on the path of the process.
Entropy can be defined as a function of the basic state parameters:

S = f 1 (P,V) ; S = f 2 (P,T) ; S = f 3 (V,T) ; (3.3)

or for specific entropy:

s = f 1 (P,υ) ; s = f 2 (P,T) ; S = f 3 (υ,T) ; (3.4)

Since entropy does not depend on the type of process and is determined by the initial and final states of the working fluid, only its change in a given process is found, which can be found using the following equations:

Ds = c v ln(T 2 /T 1) + R ln(υ 2 /υ 1); (3.5)
Ds = c p ln(T 2 /T 1) - R ln(P 2 /P 1) ; (3.6)
Ds = c v ln(P 2 /P 1) + c p ln(υ 2 /υ 1) . (3.7)

If the entropy of the system increases (Ds > 0), then heat is supplied to the system.
If the entropy of the system decreases (Ds< 0), то системе отводится тепло.
If the entropy of the system does not change (Ds = 0, s = Const), then heat is not supplied or removed to the system (adiabatic process).

Carnot cycle and theorems.

The Carnot cycle is a circular cycle consisting of 2 isothermal and 2 adiabatic processes. The reversible Carnot cycle in p,υ- and T,s-diagrams is shown in Fig. 3.1.

1-2 – reversible adiabatic expansion at s 1 = Const. The temperature decreases from T 1 to T 2.
2-3 – isothermal compression, heat removal q 2 to a cold source from the working fluid.
3-4 – reversible adiabatic compression at s 2 =Const. The temperature rises from T 3 to T 4.
4-1 – isothermal expansion, supply of heat q 1 to the hot source to the working fluid.
The main characteristic of any cycle is thermal efficiency(t.k.p.d.).

h t = L c / Q c, (3.8)

h t = (Q 1 – Q 2) / Q 1.

For a reversible Carnot cycle t.k.p.d. determined by the formula:

h tk = (T 1 – T 2) / T 1. (3.9)

this implies Carnot's 1st theorem :
|| "The thermal efficiency of a reversible Carnot cycle does not depend on
|| properties of the working fluid and is determined only by temperatures
|| sources."
From a comparison of an arbitrary reversible cycle and a Carnot cycle it follows Carnot's 2nd theorem:
|| "The reversible Carnot cycle is the best cycle in || a given temperature range"
Those. t.k.p.d. The Carnot cycle is always greater than the coefficient of efficiency. arbitrary loop:
h tк > h t . (3.10)

Topic 4. Thermodynamic processes.

Gas is one of four states of aggregation in which a substance can exist.

The particles that make up the gas are very mobile. They move almost freely and chaotically, periodically colliding with each other like billiard balls. Such a collision is called elastic collision . During a collision, they dramatically change the nature of their movement.

Since in gaseous substances the distance between molecules, atoms and ions is much greater than their sizes, these particles interact very weakly with each other, and their potential interaction energy is very small compared to the kinetic energy.

The connections between molecules in a real gas are complex. Therefore, it is also quite difficult to describe the dependence of its temperature, pressure, volume on the properties of the molecules themselves, their quantity, and the speed of their movement. But the task is greatly simplified if, instead of real gas, we consider its mathematical model - ideal gas .

It is assumed that in the ideal gas model there are no attractive or repulsive forces between molecules. They all move independently of each other. And the laws of classical Newtonian mechanics can be applied to each of them. And they interact with each other only during elastic collisions. The time of the collision itself is very short compared to the time between collisions.

Classical ideal gas

Let's try to imagine the molecules of an ideal gas as small balls located in a huge cube at a great distance from each other. Because of this distance, they cannot interact with each other. Therefore, their potential energy is zero. But these balls move at great speed. This means they have kinetic energy. When they collide with each other and with the walls of the cube, they behave like balls, that is, they bounce elastically. At the same time, they change the direction of their movement, but do not change their speed. This is roughly what the motion of molecules in an ideal gas looks like.

1. The potential energy of interaction between molecules of an ideal gas is so small that it is neglected compared to kinetic energy.
2. Molecules in an ideal gas are also so small that they can be considered material points. And this means that they total volume is also negligible compared to the volume of the vessel in which the gas is located. And this volume is also neglected.
3. The average time between collisions of molecules is much greater than the time of their interaction during a collision. Therefore, the interaction time is also neglected.

Gas always takes the shape of the container in which it is located. Moving particles collide with each other and with the walls of the container. During an impact, each molecule exerts some force on the wall for a very short period of time. This is how it arises pressure . The total gas pressure is the sum of the pressures of all molecules.

Ideal gas equation of state

The state of an ideal gas is characterized by three parameters: pressure, volume And temperature. The relationship between them is described by the equation:

Where R - pressure,

V M - molar volume,

R - universal gas constant,

T - absolute temperature (degrees Kelvin).

Because V M = V / n , Where V - volume, n - the amount of substance, and n = m/M , That

Where m - gas mass, M - molar mass. This equation is called Mendeleev-Clayperon equation .

At constant mass the equation becomes:

This equation is called united gas law .

Using the Mendeleev-Cliperon law, one of the gas parameters can be determined if the other two are known.

Isoprocesses

Using the equation of the unified gas law, it is possible to study processes in which the mass of a gas and one of the most important parameters - pressure, temperature or volume - remain constant. In physics such processes are called isoprocesses .

From The unified gas law leads to other important gas laws: Boyle-Mariotte law, Gay-Lussac's law, Charles's law, or Gay-Lussac's second law.

Isothermal process

A process in which pressure or volume changes but temperature remains constant is called isothermal process .

In an isothermal process T = const, m = const .

The behavior of a gas in an isothermal process is described by Boyle-Mariotte law . This law was discovered experimentally English physicist Robert Boyle in 1662 and French physicist Edme Mariotte in 1679. Moreover, they did this independently of each other. The Boyle-Mariotte law is formulated as follows: In an ideal gas at a constant temperature, the product of the gas pressure and its volume is also constant.

The Boyle-Marriott equation can be derived from the unified gas law. Substituting into the formula T = const , we get

p · V = const

That's what it is Boyle-Mariotte law . From the formula it is clear that the pressure of a gas at constant temperature is inversely proportional to its volume. The higher the pressure, the lower the volume, and vice versa.

How to explain this phenomenon? Why does the pressure of a gas decrease as the volume of a gas increases?

Since the temperature of the gas does not change, the frequency of collisions of molecules with the walls of the vessel does not change. If the volume increases, the concentration of molecules becomes less. Consequently, per unit area there will be fewer molecules that collide with the walls per unit time. The pressure drops. As the volume decreases, the number of collisions, on the contrary, increases. Accordingly, the pressure increases.

Graphically, an isothermal process is displayed on a curve plane, which is called isotherm . She has a shape hyperboles.

Each temperature value has its own isotherm. The higher the temperature, the higher the corresponding isotherm is located.

Isobaric process

The processes of changing the temperature and volume of a gas at constant pressure are called isobaric . For this process m = const, P = const.

The dependence of the volume of a gas on its temperature at constant pressure was also established experimentally French chemist and physicist Joseph Louis Gay-Lussac, who published it in 1802. That is why it is called Gay-Lussac's law : " Etc and constant pressure, the ratio of the volume of a constant mass of gas to its absolute temperature is a constant value."

At P = const the equation of the unified gas law turns into Gay-Lussac equation .

An example of an isobaric process is a gas located inside a cylinder in which a piston moves. As the temperature rises, the frequency of molecules hitting the walls increases. The pressure increases and the piston rises. As a result, the volume occupied by the gas in the cylinder increases.

Graphically, an isobaric process is represented by a straight line, which is called isobar .

The higher the pressure in the gas, the lower the corresponding isobar is located on the graph.

Isochoric process

Isochoric, or isochoric, is the process of changing the pressure and temperature of an ideal gas at constant volume.

For an isochoric process m = const, V = const.

It is very simple to imagine such a process. It occurs in a vessel of a fixed volume. For example, in a cylinder, the piston in which does not move, but is rigidly fixed.

The isochoric process is described Charles's law : « For a given mass of gas at constant volume, its pressure is proportional to temperature" The French inventor and scientist Jacques Alexandre César Charles established this relationship through experiments in 1787. In 1802, it was clarified by Gay-Lussac. Therefore this law is sometimes called Gay-Lussac's second law.

At V = const from the equation of the unified gas law we get the equation Charles's law or Gay-Lussac's second law .

At constant volume, the pressure of a gas increases if its temperature increases. .

On graphs, an isochoric process is represented by a line called isochore .

The larger the volume occupied by the gas, the lower the isochore corresponding to this volume is located.

In reality, no gas parameter can be maintained unchanged. This can only be done in laboratory conditions.

Of course, an ideal gas does not exist in nature. But in real rarefied gases at very low temperatures and pressures no higher than 200 atmospheres, the distance between the molecules is much greater than their sizes. Therefore, their properties approach those of an ideal gas.

>>Physics and Astronomy >>Physics 10th grade >>Physics: Equation of state of an ideal gas

Ideal gas state

We will devote today's physics lesson to the topic of the equation of state of an ideal gas. However, first, let's try to understand such a concept as the state of an ideal gas. We know that particles of real existing gases, such as atoms and molecules, have their own sizes and naturally fill some volume in space, and accordingly they are slightly dependent on each other.

When interacting between gas particles, physical forces burden their movement and thereby limit their maneuverability. Therefore, gas laws and their consequences, as a rule, are not violated only for rarefied real gases. That is, for gases, the distance between the particles of which significantly exceeds the intrinsic size of the gas particles. In addition, the interaction between such particles is usually minimal.

Therefore, gas laws at natural atmospheric pressure have an approximate value, and if this pressure is high, then the laws do not apply.

Therefore, in physics it is customary to consider such a concept as the state of an ideal gas. Under such circumstances, particles are usually regarded as certain geometric points that have microscopic dimensions and do not have any interaction with each other.

Ideal gas equation of state

But the equation that connects these microscopic parameters and determines the state of the gas is usually called the equation of state of an ideal gas.

Such zero parameters, without which it is impossible to determine the state of the gas, are:

The first parameter includes pressure, which is designated by the symbol - P;
The second parameter is volume –V;
And the third parameter is temperature – T.
From the previous section of our lesson, we already know that gases can act as reactants or be products in chemical reactions, therefore, under normal conditions, it is difficult to make gases react with each other, and for this it is necessary to be able to determine the number of moles of gases under conditions that different from normal.

But for these purposes they use the equation of state of an ideal gas. This equation is also commonly called the Clapeyron-Mendeleev equation.

Such an equation of state for an ideal gas can be easily obtained from the formula for the dependence of pressure and temperature, describing the gas concentration in this formula.

This equation is called the ideal gas equation of state.

n is the number of moles of gas;
P – gas pressure, Pa;
V – gas volume, m3;
T – absolute gas temperature, K;
R – universal gas constant 8.314 J/mol×K.

For the first time, an equation that helps establish the relationship between pressure, volume and temperature of gases was obtained and formulated in 1834 by the famous French physicist Benoit Clapeyron, who worked for a long time in St. Petersburg. But Dmitry Ivanovich Mendeleev, the great Russian scientist, first used it in 1874, but before that he obtained the formula by combining Avogadro’s law with the law that Clapeyron formulated.

Therefore, the law that allows us to draw conclusions about the nature of the behavior of gases was commonly called the Mendeleev-Clapeyron law in Europe.

Also, you should pay attention to the fact that when the volume of gas is expressed in liters, the Clapeyron-Mendeleev equation will have the following form:

I hope that you did not have any problems studying this topic and now you have an idea of ​​​​what the equation of state of an ideal gas is and you know that with its help you can calculate the parameters of real gases in the case when the physical conditions of the gases are close to normal conditions.

The state of gases is characterized by pressure R, temperature 7, and volume V. The relationship between these quantities is determined by the laws of the gas state.

Oil and natural gases have significant deviations from the laws of ideal gases due to the interaction between molecules, which occurs when real gases are compressed. The degree of deviation of the compressibility of real gases from ideal ones is characterized by the compressibility coefficient z, which shows the ratio of the volume of a real gas to the volume of an ideal gas under the same conditions.

In a reservoir, hydrocarbon gases can be found in a variety of conditions. With an increase in pressure from O to 3-4 MPa, the volume of gases decreases. In this case, the hydrocarbon gas molecules come closer together and the attractive forces between them help the external forces that compress the gas. When a hydrocarbon gas is highly compressed, the intermolecular distances are so small that repulsive forces begin to resist further reduction in volume and the compressibility of the gas decreases.

In practice, the state of real hydrocarbon gases at various temperatures and pressures can be described based on the Clapeyron equation:

P-V=z-m-R-T (2.9)

Where R - pressure gz. Pa; V" - volume occupied by gas at a given pressure, m 3 ; T - gas mass, kg; R- gas constant, J/(kg-K); T- temperature, K; G - compressibility factor.

The compressibility coefficient is determined from graphs constructed from experimental data.

State of hydrocarbon gas-liquid systems with changes in pressure and temperature.

When oil and gas move in the formation, wellbore, collection and treatment systems, pressure and temperature change, which causes a change in the phase state of hydrocarbons - a transition from liquid to gaseous state and vice versa. Since oil and gas consist of a large number of components with different properties, under certain conditions some of these components can be in the liquid phase, and the other in the vapor (gas) phase. It is obvious that the patterns of movement of a single-phase system in the formation and wellbore are significantly different from the patterns of multiphase movement. Conditions for long-distance transport of oil and gas and subsequent processing require the separation of easily evaporating components from the liquid condensed fraction. Therefore, the choice of field development technology and in-field oil and gas treatment system is largely related to the study of the phase state of hydrocarbons under changing thermodynamic conditions.

Phase transformations of hydrocarbon systems are illustrated by phase diagrams showing the relationship between pressure, temperature and specific volume of a substance.

In Fig. 2.2, A The state diagram of pure gas (ethane) is shown. The solid lines in the diagram show the relationship between pressure and the specific volume of a substance at constant temperatures. The lines passing through the area bounded by the dotted curve have three characteristic sections. If we consider one of the lines of the high pressure region, then at first the increase in pressure is accompanied by a slight increase in the specific volume of the substance, which is compressible and in this region is in a liquid state.

Rice. 2.2. Pure gas phase diagram

At a certain pressure, the isotherm breaks sharply and looks like a horizontal line. At constant pressure there is a continuous increase in the volume of the substance. In this area, the liquid evaporates and enters the vapor phase. Evaporation ends at the point of the second break of the isotherm, after which the change in volume is accompanied by an almost proportional decrease in pressure. In this region, all matter is in gaseous form.

state (in the vapor phase). The dotted line connecting the break points of the isotherms limits the region of transition of a substance from a liquid to a vapor state or vice versa (in the direction of decreasing specific volumes). This region corresponds to the conditions under which a substance is simultaneously in two states, liquid and gaseous (region of a two-phase state of a substance). The dotted line located to the left of point C is called vaporization point curve. The coordinates of the points on this line are the pressure and temperature at which the substance begins to boil. To the right of point C lies a dotted line called dew point curve or dew points. It shows at what pressures and temperatures vapor condensation begins - the transition of a substance into a liquid state. Point C lying at the top of the two-phase region is called critical point. At the pressure and temperature corresponding to this point, the properties of the vapor and liquid phases are the same. In addition, for a pure substance, the critical point determines the highest values ​​of pressure and temperature at which the substance can simultaneously be in a two-phase state. When considering an isotherm that does not cross the two-phase region, it is clear that the properties of the substance change continuously and the transition of the substance from the liquid to the gaseous state or vice versa occurs without passing through the two-phase state.

In Fig. 2.2, b The state diagram of ethane is shown, rearranged in pressure-temperature coordinates. Since a pure substance passes from one phase state to another at constant pressure, the curves of the evaporation and condensation points in this diagram coincide and end with the critical point C. The resulting line delimits the regions of liquid and vaporous substances. A substance can be in a two-phase state only at pressures and temperatures corresponding to the coordinates of this line.